Thermodynamics of Neptunium(V) Fluoride and Sulfate at Elevated Temperatures

Linfeng Rao,1 Guoxin Tian,1 Yuanxian Xia2 and Judah I. Friese2

1Lawrence Berkeley National Laboratory, Berkeley, CA 94720 2Pacific Northwest National Laboratory, Richland, WA 99352

Abstract – Complexation of neptunium(V) with fluoride and sulfate at elevated temperatures was studied by microcalorimetry. Thermodynamic parameters, including the equilibrium constants and enthalpy of protonation of fluoride and sulfate, and the enthalpy of complexation between Np(V) and fluoride and sulfate at 25 – 70oC were determined. Results show that the complexation of Np(V) with fluoride and sulfate is endothermic and that the complexation is enhanced by the - increase in temperature – a three-fold increase in the stability constants of NpO2F(aq) and NpO2SO4 as the temperature is increased from 25 to 70oC.

I. INTRODUCTION characteristic absorption peak in the near IR region and the concentration of Np(V) was determined by the Neptunium is one of the radionuclides of concern in absorbance at 980 nm using the molar absorption the post-closure chemical environment in the proposed coefficient of 395 M-1cm-1. Gran’s potentiometric method Yucca Mountain repository because of its mobility and [4] was used to determine the concentration of perchloric long half-life (2.14 × 106 years). It is likely that 237Np, 129I acid in the Np(V) stock solution. Solutions of fluoride and 99 and Tc will be the major contributors to the potential sulfate were prepared by dissolving solid NaF or Na2SO4 total annual dose from the repository beyond 10,000 years in water. The ionic strength of all the working solutions in [1]. this study was maintained at 1.0 M (25°C) by adding appropriate amounts of sodium perchlorate as the The postclosure chemical environment is expected to background electrolyte. be at near neutral pH, slightly oxidizing, and at elevated

temperatures for thousands of years [1]. Np(V) is II.B. Microcalorimetry expected to be the most stable oxidation state under these

conditions. Its complexation with inorganic ligands that Calorimetric titrations were performed at 25, 40, 55 could be present, such as OH-, F-, SO 2- and CO 2-, plays 4 3 and 70oC to determine the equilibrium constants and the an important role in determining its migration in the enthalpy of protonation of fluoride and sulfate, and the repository. To predict the migration behavior of enthalpy of complexation of Np(V) with fluoride and neptunium, it is necessary to have sufficient and reliable sulfate. The titrations were conducted on an isothermal thermodynamic data on its complexation at elevated microcalorimeter (Model ITC 4200, Calorimetry Science temperatures. However, such data are scarce and scattered Corp.). The performance of the calorimeter has been for 25oC, and nonexistent for elevated temperatures [2]. tested by measuring the enthalpy of protonation of To provide reliable thermodynamic data, we have started tris(hydroxymethyl)-aminomethane (THAM). The result investigations of the complexation of actinides at elevated was –(47.7 ± 0.2) kJ·mol-1 at 25ºC, in excellent agreement temperatures. Thermodynamic parameters, including with the value in the literature [3]. Details of the formation constants, enthalpy and heat capacity of microcalorimeter are provided elsewhere [5]. complexation are experimentally determined. This paper summarizes the results of the complexation of Np(V) with In the titrations of the protonation of fluoride and fluoride and sulfate at 25 - 70oC. sulfate, the initial cell solutions (~ 0.90 cm3 at 25oC) containing fluoride or sulfate were titrated with HClO4 in II. EXPERIMENTAL increments of 5 µl. In the titrations of the complexation of Np(V), the initial cell solutions (~ 0.90 cm3 at 25oC) II.A. Chemicals containing Np(V) were titrated with solutions of fluoride or sulfate in increments of 5 µl. Multiple titrations were All chemicals were reagent grade or higher. Water conducted for each ligand at each temperature. For each from a Milli-Q system was used in preparations of all the titration, n additions were made (usually n = 40 - 50), solutions. The stock solution of Np(V) in perchloric acid resulting in n experimental values of the heat generated in was prepared based on procedures previously reported the reaction cell (Qex,j, where j = 1 to n). These values [3]. The oxidation state of Np(V) was confirmed by the were corrected for the heat of dilution of the titrant (Qdil,j), which was determined in separate runs. The net reaction heat at the j-th point (Qr,j) was obtained from the difference: Qr,j = Qex,j - Qdil,j. The program Letagrop [6] was used to analyze the data and calculate the thermodynamic parameters.

III. RESULTS AND DISCUSSION

III.A. Protonation of Fluoride and Sulfate at Elevated Temperatures

Prior to the determination of the thermodynamic parameters of neptunium complexation with fluoride and sulfate at elevated temperatures, the protonation constants and enthalpy of protonation of the ligands at these temperatures must be known or, if not available, must be determined. The frequently used potentiometry with a glass electrode is not suitable for such studies because of two reasons: 1) corrosion of the glass electrode by HF, especially at elevated temperatures, precludes the use of glass pH electrodes; 2) the pK of HSO4- (~ 1) is too low a Fig. 1. Calorimetric titrations of the protonation of to be accurately measured by acid-base potentiometry. On fluoride at different temperatures. Initial cup solution: the other hand, titration calorimetry is ideal for such 0.90 mL, C = 0.0333 M, titrant: 0.1032 M HClO , 5 measurements. Because the experimentally observed heat NaF 4 depends on the enthalpy of reaction as well as the changes µL/addition; I = 1 M NaClO4. in speciation during the titration, it is possible to obtain the equilibrium constant(s) and the enthalpy of protonation simultaneously by fitting the reaction heat.

Figure 1 shows the reaction heat of protonation of fluoride determined by calorimetry at different temperatures. Negative values of heat indicate that the overall reaction is endothermic. From the data of multiple calorimetric titrations at each temperature, the equilibrium constants and enthalpy of protonation of fluoride and sulfate at 25 - 70oC are calculated and shown in Figure 2 and Table 1. The trends can be summarized as follows: 1) protonation constants of fluoride and sulfate increase as the temperature is elevated; 2) protonation of fluoride and sulfate is endothermic and becomes more endothermic at higher temperatures. From the linear correlation between the enthalpy and temperature (Figure 2), the heat capacity of protonation of fluoride and sulfate was calculated to be (96 ± 13) J·K-1·mol-1 and (599 ± 16) J·K-1·mol-1, respectively. It is worth noting that, as Table 1 shows, the equilibrium constant and the enthalpy of protonation of o sulfate at 25 C and I = 1.0 M NaClO4 from this work are similar to the calculated values by Dickson et al. at 25oC and I = 1.0 m NaCl [7].

Fig. 2. Effect of temperature on the protonation constant and enthalpy (in kJ·mol-1) of fluoride and sulfate. I = 1 M NaClO4.

TABLE 1. Thermodynamic parameters for the III.B. Complexation of Np(V) with Fluoride and protonation and complexation of fluoride and sulfate with Sulfate at Elevated Temperatures Np(V). A representative calorimetric titration of the 0 Reaction t log10K log10K ∆H complexation of Np(V) with fluoride is shown in Figure o C I = 1.0 M I = 0 kJ·mol-1 3, as the total reaction heat and Np(V) speciation vs. the NaClO4 volume of titrant. H+ + F- = 25 2.65±0.03 2.87 12.9±0.3 As previously discussed, the observed reaction heat is HF(aq) a function of a number of parameters, including the 40 2.78±0.03 3.01 13.1±0.6 equilibrium constant and enthalpy of protonation of the ligands (fluoride and sulfate), the concentrations of 55 3.20±0.06 3.44 16.0±0.6 reactants (Np(V), ligands and proton), and the enthalpy of 70 3.30±0.15 3.55 17.2±0.6 complexation as well as the stability constants of the + 2- complexes that form in the titration. To calculate the H + SO4 = 25 1.07±0.09 1.84 22.7±3.0 - enthalpy of complexation from the reaction heat, the other HSO4 1.06±0.02a 19.9±0.8a parameters should be known. In this work, we have used the protonation constants and enthalpy of fluoride and 40 1.14±0.12 1.92 32.0±2.0 sulfate determined by calorimetry and the stability - constants of NpO2F(aq) and NpO2SO4 from ref [8]. 55 1.28±0.09 2.08 40.0±5.0 Because some of the stability constants of NpO2F(aq) and - 70 1.38±0.09 2.21 50.0±5.0 NpO2SO4 were determined in a slightly different temperature range (24 – 60oC) [8], interpolation and + - NpO2 + F = 24 1.42±0.10 extrapolation were performed to obtain the stability o NpO2F(aq) constants at 25, 40, 55 and 70 C. These constants and the 25 b 1.58 1.47±0.10 18.0±0.3 enthalpy of complexation are shown in Figure 4, and 35 1.63±0.03 summarized in Table 1 together with the parameters obtained for the protonation of fluoride and sulfate. b 40 1.63±0.10 1.75 12.2±0.6

51 1.77±0.04

55 1.78±0.10b 1.91 9.4±0.8 60 1.80±0.03 70 1.94±0.10b 2.08 8.2±1.2 + 2- NpO2 + SO4 = 25 0.49±0.31 0.97 23.1±2.5 - NpO2SO4 40 0.65±0.19 1.14 20.0±1.6 50 0.73±0.16 55 0.78±0.20b 1.29 25.1±0.3 70 0.93±0.20b 1.47 30.4±1.3

a From Ref.[7], I = 1.0 m NaCl. b Extrapolated or interpolated stability constants of - NpO2F(aq) and NpO2SO4 from data in Ref.[8]. The error limits (3σ) of the stability constants have been enlarged Fig. 3. Calorimetric titration of Np(V) fluoride by two times to account for the uncertainties that might o complexation (I = 1.0 M NaClO4, t = 25 C). Cup: 0.900 have been introduced by the interpolation or -3 -7 mL CNp = 1.78 × 10 M, CH = 3 × 10 M; titrant: 1 M extrapolation. NaF, 5 µL/addition.

interaction parameters at 25oC (in kg mol-1), taken from + - Ref.[9], include the following: ε(H , ClO4 ) = 0.14 ± 0.02, + - + 2- ε(Na , F ) = 0.02 ± 0.02, ε(Na , SO4 ) = -0.12 ± 0.02, + - + - ε(Na , HSO4 ) = -0.01 ± 0.02 and ε(NpO2 , ClO4 ) = 0.25 + - ± 0.05. ε(Na , NpO2SO4 ) is not available, but we + - assumed that it is equal to ε(Na , NpO2CO3 ) = -0.18 ± 0.15 kg mol-1 [9]. For the calculation of log K0 at temperatures other than 25oC, the Debye-Huckel term in Eq. (1) was calculated with the values of A at different temperatures [9] and the ion interaction parameters at 25oC were used because the values at other temperatures were not known. Using the interaction parameters at 25oC may introduce some errors into the log K0 at other temperatures. However, the errors are probably small, -1 since the values of (∂ε/∂T)p are usually ≤ 0.005 kg mol K-1 for temperatures below 200oC [9]. Besides, the values

of (∂ε/∂T)p for the reactants and products may balance out each other so that the ∆ε for many reactions remains approximately constant up to 100oC [10]. The calculated values of log K0 at different temperatures are listed in

Table 1. Fig. 4. Effect of temperature on the stability constants of

NpO F(aq) and NpO SO - [8] and the enthalpy of 2 2 4 ACKNOWLEDGMENTS complexation (in kJ·mol-1). I = 1 M NaClO . 4

This work was supported by the Director, OST&I Program, Office of Civilian Radioactive Waste Data in Figure 4 indicate that the complexation of Management, U. S. Department of Energy, under Np(V) with fluoride and sulfate is enhanced at higher Contract No. DE-AC02-05CH11231 at Lawrence temperatures. The stability constants of the complexes, Berkeley National Laboratory. NpO F(aq) and NpO SO - increase by about 3 times as 2 2 4 the temperature is elevated from 25 to 70oC. The - REFERENCES enthalpies of complexation for NpO2F(aq) and NpO2SO4 are both endothermic, but the trends with the change in [1] OCRWM; “Yucca Mountain Science and temperature are opposite. As the temperature is increased, Engineering Report Rev.1”, DOE/RW-0539-1, the enthalpy of complexation for NpO2F(aq) decreases - Office of Civilian Radioactive Waste while that for NpO2SO4 increases. Assuming linear Management: North Las Vegas, NV. 2002. correlations between ∆H and t, the heat capacity of - [2] R. Guillaumont, T. Fanghanel, J. Fuger, I. complexation for NpO2F(aq) and NpO2SO4 was Grenthe, V. Neck, D. A. Palmer, M. H. Rand; calculated to be -(267 ± 20) J·K-1·mol-1 and (244 ± 51) -1 -1 “Update on the chemical thermodynamics of J·K ·mol , respectively. uranium, neptunium, plutonium, americium and It is desirable to extrapolate the equilibrium constants technetium”, (Mompean, F. J., Illemassene, M., at I = 1.0 M to the values in an infinite dilute solution (I = Domenech-Orti, C., Ben Said, K., eds.), 0). We have used the SIT (Specific Ion Interaction) Amsterdam: Elsevier B.V., 2003. approach described in the literature [9] to calculate the [3] L. Rao, T. G. Srinivasan, A. Yu. Garnov, P. equilibrium constants at I = 0. For the reactions shown in Zanonato, P. Di Bernardo, A. Bismondo; 0 Hydrolysis of Neptunium(V) at Variable Table 1, the equilibrium constants at I = 0 (log K ) are o related to log K at other ionic strengths by the following Temperatures (10 - 85 C), Geochim. equation: Cosmochim. Acta, 68, 4821-4830 (2004). [4] G. Gran; Analyst, 77, 661 (1952). 2 0 [5] P. Zanonato, P. Di Bernardo, A. Bismondo, G. log K – ∆Z × D = log K – ∆εIm (1) Liu, X. Chen, L. Rao; Hydrolysis of 2 2 2 Uranium(VI) at Variable Temperatures (10 - 85 where ∆Z = {Σ(Z products) - Σ(Z reactants)}, D is the Debye- o 1/2 C), J. Am. Chem, Soc. 126, 5515-5522 (2004). Huckel term used in the SIT method and D = AIm /(1 + 1/2 [6] R. Arnek; Arkiv Kemi, 32, 81 (1970). 1.5AIm ), Im is the ionic strength in molality, and ε is the ion interaction parameter used in the SIT method [9]. The [7] A. G. Dickson, D. J. Wesolowski, D. A. Palmer, R. E. Mesmer; Dissociation constant of Bisulfate Ion in Aqueous Sodium Chloride Solutions to 250oC. J. Phys. Chem., 94, 7978-7985 (1990). [8] Y. Xia, J. I. Friese, D. A. Moore, L. Rao; Stability Constants of Np(V) Complexes with Fluoride and Sulfate at Variable Temperatures. J. Radioanal. Nucl. Chem., 268, 3 (2006). [9] W. Hummel, G. Anderegg, I. Puigdomènech, L. Rao, O. Tochiyama; “Chemical Thermodynamics of Compounds and Complexes of: U, Np, Pu, Am, Tc, Zr, Ni and Se with Selected Organic Ligands”, (Mompean, F. J.; Illemassene, M.; Perrone, J. eds.), Amsterdam: Elsevier B.V. (2005). [10] A. V. Plyasunov, I. Grenthe; Geochim. Cosmochim. Acta, 58, 3561-3582 (1994).